For the reaction:
$I ^{-}+ ClO _{3}^{-}+ H _{2} SO _{4} \longrightarrow Cl ^{-}+ HSO _{4}^{-}+ I _{2}$
The correct statement(s) in the balanced equation is/are:
- stoichiometric coefficient of $HSO_4^-$ is $6$
- iodide is oxidised
- sulphur is reduced
- $H_2O$ is one of the products
The Correct Option is D
Solution and Explanation
$ClO _{3}^{-}+6 I ^{-}+6 H _{2} SO _{4} \longrightarrow 3 I _{2}+ Cl ^{-}+6 HSO _{4}^{-}+3 H _{2} O$
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Concepts Used:
Redox Reactions
Redox Reaction:
Redox reactions are chemical reactions where oxidation and reduction take place simultaneously. In this type of reaction, there is a gain of electrons for one chemical species while the other loses electrons or simply involves transfer of electrons. The species that loses electrons is oxidized while the one that gains electrons is reduced.
Types of Redox Reactions:
Redox reactions can be differentiated into 4 categories namely combination reactions, decomposition reactions, displacement reactions, and disproportionation reactions. Each is explained separately below:
Combination Reaction:
In this, the molecules combine to form new compounds. For example, when magnesium reacts to nitrogen.
Decomposition Reaction:
Opposite to the combination reaction, here there is a breakdown of compounds to simpler substances. For example, electrolysis of water.
Displacement Reaction:
In this, the more reactive metal will displace the less reactive one in a chemical reaction. The reactivity of an element is represented in a series called the reactivity series (arranged in decreasing order of reactivity) which makes it easier to determine the chemical reaction and its products.
Disproportionation Reaction:
This is a peculiar type of reaction where an element showing a particular oxidation state will be oxidized and reduced simultaneously. Another thing to note is that these reactions will always have an element that can exhibit three oxidation states.